Ionic and Covalent Bonding Describing Ionic Bonds n

Ionic and Covalent Bonding

Describing Ionic Bonds n An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions. – This type of bond involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal). – The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” with a noble gas. 2

Describing Ionic Bonds n Such noble gas configurations and the corresponding ions are particularly stable. – The atom that loses the electron becomes a cation (positive). – The atom that gains the electron becomes an anion (negative). 3

Describing Ionic Bonds n Consider the transfer of valence electrons from a sodium atom to a chlorine atom. e- ions are electrostatically attracted to – The resulting one another. – The attraction of these oppositely charged ions for one another is the ionic bond. 4

Lewis Electron-Dot Symbols n A Lewis electron-dot symbol is a symbol in which the electrons in the valence shell of an atom or ion are represented by dots placed around the letter symbol of the element. . P. : S. : Cl. : Ar: Na. . Mg. . Al. . Si. . . Group II Group IV Group VIII : : : : . Group III : : Group I – Note that the group number indicates the number of valence electrons. 5

Lewis Electron-Dot Formulas n A Lewis electron-dot formula is an illustration used to represent the transfer of electrons during the formation of an ionic bond. – As an example, let’s look at the transfer of electrons from magnesium to fluorine to form magnesium fluoride. 6

Lewis Electron-Dot Formulas • The magnesium has two electrons to give, whereas the fluorines have only one “vacancy” each. : : 2+ Mg [: F: ] : : n : : [: F: ] - . F: : : . Mg. : F. - Consequently, magnesium can accommodate two fluorine atoms. 7

Energy Involved in Ionic Bonding n The transfer of an electron from a sodium atom to a chlorine atom is not in itself energetically favorable; it requires 147 k. J/mol of energy. – However, 493 k. J of energy is released when these oppositely charged ions come together. – An additional 293 k. J of energy is released when the ion pairs solidify. – This “lattice energy” is the negative of the energy released when gaseous ions form an ionic solid. The next slide illustrates this. (See Animation: Born Haber Cycle) 8

Energy Involved in Ionic Bonding 9

Electron Configurations of Ions n As metals lose electrons to form cations and establish a “noble gas” configuration, the electrons are lost from the valence shell first. – For example, magnesium generally loses two electrons from its 3 s subshell to look like neon. [Ne]3 s 2 [Ne] 10

Electron Configurations of Ions n Transition metals also lose electrons from the valence shell first, which is not the last subshell to fill according to the aufbau sequence. – For example, zinc generally loses two electrons from its 4 s subshell to adopt a “pseudo”-noble gas configuration. [Ar]4 s 23 d 10 [Ar]3 d 10 11

Ionic Radii n The ionic radius is a measure of the size of the spherical region around the nucleus of an ion within which the electrons are most likely to be found. – Ionic radii increase down any column because of the addition of electron shells. – In general, across any period the cations decrease in radius. When you reach the anions, there is an abrupt increase in radius, and then the radius again decreases. 12

Ionic Radii 13

Ionic Radii n Within an isoelectronic group of ions, the one with the greatest nuclear charge will be the smallest. – For example, look at the ions listed below. All have 18 electrons – Note that they all have the same number of electrons, but different numbers of protons. 14

Ionic Radii n In this group, calcium has the greatest nuclear charge and is, therefore, the smallest. All have 18 electrons – Sulfur has only 16 protons to attract its 18 electrons and, therefore, has the largest radius. 15

Covalent Bonds n When two nonmetals bond, they often share electrons since they have similar attractions for them. This sharing of valence electrons is called the covalent bond. – These atoms will share sufficient numbers of electrons in order to achieve a noble gas electron configuration (that is, eight valence electrons). 16

Covalent Bonds n The tendency of atoms in a molecule to have eight electrons in their outer shell (two for hydrogen) is called the octet rule. – Figure 9. 10 illustrates how two electrons can be shared by bonded hydrogen atoms. – Figure 9. 11 shows the potential energy of the atoms for various distances between nuclei. The decrease in energy is a reflection of the bonding of the atoms. 17

Lewis Structures n You can represent the formation of the covalent bond in H 2 as follows: H . +. H : H H – This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots. 18

Lewis Structures n The shared electrons in H 2 spend part of the time in the region around each atom. : H H – In this sense, each atom in H 2 has a helium configuration. 19

Lewis Structures n The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way. : : . +. Cl: : : H H Cl – Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar). 20

Lewis Structures n Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures. bonding pair : : H Cl lone pair – An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared). 21

Coordinate Covalent Bonds n When bonds form between atoms that both donate an electron, you have: A . +. B : A B – It is, however, possible that both electrons are donated by one of the atoms. This is called a coordinate covalent bond. A + : B : A B 22

Multiple Bonds In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared. – It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms. : H : : C : C H : n H or 23

Multiple Bonds Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms. C : : : : H C : n H or 24

Polar Covalent Bonds n A polar covalent bond is one in which the bonding electrons spend more time near one of the two atoms involved. – When the atoms are alike, as in the H-H bond of H 2 , the bonding electrons are shared equally (a nonpolar covalent bond). – When the two atoms are of different elements, the bonding electrons need not be shared equally, resulting in a “polar” bond. 25

Polar Covalent Bonds n d : : d- : : For example, the bond between carbon and oxygen in CO 2 is considered polar because the shared electrons spend more time orbiting the oxygen atoms. + d- – The result is a partial negative charge on the oxygens (denoted d-) and a partial positive charge on the carbon (denoted d+) 26

Polar Covalent Bonds n Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. – In general, electronegativity increases from the lower-left corner to the upper-right corner of the periodic table. – The current electronegativity scale, developed by Linus Pauling, assigns a value of 4. 0 to fluorine and a value of 0. 7 to cesium. 27

Electronegativities 28

Polar Covalent Bonds n The absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the polarity of the bond. – When this difference is small (less than 0. 5), the bond is nonpolar. – When this difference is large (greater than 0. 5), the bond is considered polar. – If the difference exceeds approximately 1. 8, sharing of electrons is no longer possible and the bond becomes ionic. 29

Writing Lewis Dot Formulas n The Lewis electron-dot formula of a covalent compound is a simple twodimensional representation of the positions of electrons in a molecule. – Bonding electron pairs are indicated by either two dots or a dash. – In addition, these formulas show the positions of lone pairs of electrons. 30

Writing Lewis Dot Formulas n The following rules allow you to write electron-dot formulas even when the central atom does not follow the octet rule. n This will be illustrated using the molecule for nitric acid HNO 3 31

Writing Lewis Dot Formulas n Step 1: Determine the most symmetrical skeleton for the molecule. [The least electronegative element is used as the central atom. ] O OH N O 32

Writing Lewis Dot Formulas n Step 2: Determine the number of bonds. S=N-A n A =Total of valence electrons in the molecular formula. Available electrons. [group numbers] – For a polyatomic anion, add the number of negative charges to this total. – For a polyatomic cation, subtract the number of positive charges from this total. 33

Writing Lewis Dot Formulas O OH N O 3 oxygen x 6 e- = 18 e 1 nitrogen x 5 e- = 5 e 1 hydrogen x 1 e- = 1 e- 24 e- available 34

Writing Lewis Dot Formulas n Step 2: Determine the number of bonds. S=N-A n N =Total of valence electrons needed to complete outer energy level. Needed electrons. 35

Writing Lewis Dot Formulas O OH N O 3 oxygen x 8 e- = 24 e 1 nitrogen x 8 e- = 8 e 1 hydrogen x 2 e- = 2 e- 34 e- available 36

Writing Lewis Dot Formulas S=N–A S = 34 – 24 = 10 10 electrons involved in bonding Each bond contains 2 electrons. 5 bonds in this molecule 37

Writing Lewis Dot Formulas n Step 3: Distribute the electrons to the peripheral atoms first to satisfy the octet rule. O O H N O 38

Writing Lewis Dot Formulas n Step 4: Distribute any remaining electrons to the central atom. If there are fewer than eight electrons on the central atom, a multiple bond may be necessary. 39

Writing Lewis Dot Formulas O O H N O 40

Writing Lewis Dot Formulas n Try drawing Lewis dot formulas for the following covalent compound. SCl 2 S Cl : : Cl 41

Writing Lewis Dot Formulas Try drawing Lewis dot formulas for the following covalent compound. COCl 2 : O : : n C Cl: : Cl 42

Writing Lewis Dot Formulas Note that the carbon has only 6 electrons. n One of the oxygens must share a lone pair. COCl 2 : O : : n C Cl: : Cl 43

Writing Lewis Dot Formulas n Note that the carbon has only 6 electrons. n One of the oxygens must share a lone pair. COCl 2 : O : C Cl: : Cl Note that the octet rule is now obeyed. 44

Delocalized Bonding: Resonance The structure of ozone, O 3, can be represented by two different Lewis electron-dot formulas. : : O : : O: : : O O or : : n O – Experiments show, however, that both bonds are identical. 45

Delocalized Bonding: Resonance n According to theory, one pair of bonding electrons is spread over the region of all three atoms. O O O – This is called delocalized bonding, in which a bonding pair of electrons is spread over a number of atoms. 46

Delocalized Bonding: Resonance : : According to the resonance description, you describe the electron structure of molecules with delocalized bonding by drawing all of the possible electron-dot formulas. O : : O: : : O O and : : n O – These are called the resonance formulas of the molecule. 47

Exceptions to the Octet Rule n n n Although many molecules obey the octet rule, there are exceptions where the central atom has more than eight electrons. Be can produce a stable bond with only two electrons F – Be – F B can produce a stable bond with only three electrons. Cl Cl B Cl 48

Exceptions to the Octet Rule n There are exceptions where the central atom has more than eight electrons. – Generally, if a nonmetal is in the third period or greater it can accommodate as many as twelve electrons, if it is the central atom. – These elements have unfilled “d” subshells that can be used for bonding. 49

Exceptions to the Octet Rule n : F: : : P : : : : F: : : For example, the bonding in phosphorus pentafluoride, PF 5, shows ten electrons surrounding the phosphorus. F: F: : 50

Exceptions to the Octet Rule n In xenon tetrafluoride, Xe. F 4, the xenon atom must accommodate two extra lone pairs. Xe F: : : F : : : F F: : : 51

Exceptions to the Octet Rule n There are exceptions where the central atom has less than eight electrons. n n n Two small elements in the second period, also frequently break the octet rule. Be can be satisfied with only two bonds and B with only three bonds F – Be – F Cl Cl B Cl 52

Formal Charge and Lewis Structures n In certain instances, more than one feasible Lewis structure can be illustrated for a molecule. For example, H C N: or H N C: – The concept of “formal charge” can help discern which structure is the most likely. 53

Formal Charge and Lewis Structures The formal charge of an atom is determined by subtracting the number of electrons in its “domain” from its group number 1 e- 4 e- 5 e 1 e- 5 e- 4 e. H C N : or H N C : n “domain” electrons I IV V IV – The number of electrons in an atom’s “domain” is determined by counting one electron for each bond and two electrons for each lone pair. 54

Formal Charge and Lewis Structures n The most likely structure is the one with the least number of atoms carrying formal charge. If they have the same number of atoms carrying formal charge, choose the structure with the negative formal charge on the more electronegative atom. formal charge or H C 0 0 N: 0 H 0 N C: +1 -1 – In this case, the structure on the left is most likely correct. 55

Bond Length and Bond Order n Bond length (or bond distance) is the distance between the nuclei in a bond. – Knowing the bond length in a molecule can sometimes give clues as to the type of bonding present. – Covalent radii are values assigned to atoms such that the sum of the radii of atoms “A” and “B” approximate the A-B bond length. 56

Bond Length and Bond Order n Table 9. 4 lists some covalent radii for nonmetals. – For example, to predict the bond length of CCl, you add the covalent radii of the two atoms. C Cl Bond length 57

Bond Length and Bond Order n The bond order, determined by the Lewis structure, is the number of pairs of electrons in a bond. – Bond length depends on bond order. – As the bond order increases, the bond gets shorter and stronger. C C C Bond length 154 pm C 134 pm C 120 pm C Bond energy 346 k. J/mol 602 k. J/mol 835 k. J/mol 58

Bond Energy n We define the A-B bond energy (denoted BE) as the average enthalpy change for the breaking of an A-B bond in a molecule in its gas phase. – The enthalpy, DH, of a reaction is approximately equal to the sum of the bond energies of the reactants minus the sum of the bond energies of the products. – Table 9. 5 lists values of some bond energies. 59

Bond Energy n To illustrate, let’s estimate the H for the following reaction. – In this reaction, one C-H bond and one Cl-Cl bond must be broken. – In turn, one C-Cl bond and one H-Cl bond are formed. 60

Bond Energy n Referring to Table 9. 5 for the bond energies, a little simple arithmetic yields H. 61

Operational Skills n n n n n Using Lewis symbols to represent ionic bond formation. Writing electron configurations of ions. Using periodic trends to obtain relative ionic radii. Using electronegativities to obtain relative bond polarity. Writing Lewis formulas. Writing resonance structures. Using formal charges to determine the best Lewis formula. Relating bond order and bond length. Estimating DH from bond energies. 62

Animation: Born Haber Cycle (Click here to open Quick. Time animation) Return to Slide 8 63